Koç University
2010, Fall
Physical Chemistry Laboratory (Chem 301)
Experiments
Laboratory Safety Rules
ALWAYS:
· Wear safety glasses
· Wear protective clothing
· Know the location and use of all safety equipment
· Use proper techniques and procedures
· Add acid to water
· Be very cautious when testing for odors
· Use hoods whenever poisonous or irritating fumes are evolved
· Discard wastes properly - flush liquids down the sink with a large excess of water
· Report any accident, however minor, to the instructor at once
· All the times think about what you are doing
· Be alert, serious, and responsible
NEVER:
· Eat or drink in the lab
· Perform unauthorized experiments
· Leave anything unattended while it is being heated or is reacting rapidly
· Aim the opening of a test tube or flask at yourself or at anyone else
· Add water to acid
· Insert droppers, pipettes and other laboratory equipment into reagent bottles – this is a sure way of contaminating the contents
· Return unused reagents to stock bottles
· Clutter your work area
· Take unnecessary risks
· Enter chemical storage area
Experiment 1
Ideal Gas Laws
1. Purpose
The purpose of the experiment is to demonstrate the ideality of a real gas at given temperature and atmospheric pressure.
2. Introduction
An ideal gas is a gas that conforms, in physical behavior, to a particular, idealized relation between pressure, volume, and temperature called the ideal gas law. This law is a generalization containing both Boyle's law and Charles's law as special cases and states that for a specified quantity of gas, the product of the volume, V, and pressure, P, is proportional to the absolute temperature T; i.e., in equation form, PV = kT, in which k is a constant. Such a relation for a substance is called its equation of state and is sufficient to describe its macroscopic behavior.
The ideal gas law can be derived from the kinetic theory of gases and relies on the assumptions that (1) the gas consists of a large number of molecules, which are in random motion and obey Newton's laws of motion; (2) the volume of the molecules is negligibly small compared to the volume occupied by the gas; and (3) no forces act on the molecules except during elastic collisions of negligible duration.
Although no gas has these properties, the behavior of real gases is described quite closely by the ideal gas law at sufficiently high temperatures and low pressures, when relatively large distances between molecules and their high speeds overcome any interaction. A gas does not obey the equation when conditions are such that the gas, or any of the component gases in a mixture, is near its condensation point.
The ideal gas law may be written in a form applicable to any gas, according to Avogadro's law, if the constant specifying the quantity of gas is expressed in terms of the number of molecules of gas. This is done by using as the mass unit the gram-mole; i.e., the molecular weight expressed in grams. The equation of state of n gram-moles of a perfect gas can then be written as PV/T = nR, in which R is called the universal gas constant. This constant has been measured for various gases under nearly ideal conditions of high temperatures and low pressures, and it is found to have the same value for all gases: R = 8.314 J/g mol K.
According to Boyle’s law, the pressure (P) of a gas varies inversely with the volume (V) of the gas when the temperature (T) and quantity (n) are kept constant. Mathematically, this is expressed as
P ~1/V (T, n constant) Boyles’s law
By rearranging it can be found that the product of pressure and volume of a specific amount of gas is constant as long as the temperature does not change.
P1V1= constant
At the same temperature, the PV product of a gas remains constant. Therefore, it can be written as equality of the initial and the final products of pressure and volume.
P1V1=P2V2
According to Boyle’s law, volume decreases when the pressure increases, and volume increases when the pressure decreases.
When temperature rises, the kinetic energy of the molecules increases. To keep the pressure constant, the volume must expand. According to Charles’ law, the volume of a gas changes directly with the Kelvin temperature as long as the pressure and the number of the moles remains constant.
V ~ T (P and n constant) or V/T= constant
Therefore Charles’ law can be written as
2. Experimental procedure
2.1. Charles’ Law
Materials: 50 ml Erlenmeyer flask, 200 ml beaker, one-hole rubber stopper with a short piece of glass tubing inserted and attached to a piece of rubber tubing, water containers, thermometers, pinch clamps, burette clamp, hot plate, graduated cylinder, boiling chips, ice.
Procedure: Dry any moisture on the inside of the 50 ml Erlenmeyer flask. Place a one-hole rubber stopper and tubing in the neck of the flask. Set a 200 ml beaker on a hot plate. Add a few boiling chips to the bottom of the beaker. Attach the burette clamp to the neck of the flask and lower the flask into a 200 ml beaker without touching the flask of the bottom. Fasten the burette clamp to the ring stand. Pour the water into the beaker until it comes up to the neck of the flask. Leave place at the top for the water to boil without boiling over (see figure 1).
Figure 1. Setup for heating an Erlenmeyer flask in boiling water bath
Begin heating and bring the water to boil. Boil gently for 10 min to bring the temperature of the air in flask to that of the boiling water. Measure the temperature of the boiling water. Convert from degree Celsius to the corresponding Kelvin temperature.
Place the pinch clamp on the rubber tubing. In the lab there will be 3 large containers with water at different temperatures. Using the burette clamp as a holder, carefully lift the flask out of the hot water and carry it over the one of the cool water containers. Keep the stopper end of the flask (closed with pinch clamp) pointed downward, immerse the flask in the cool water (see Figure 2). Keep the flask inverted and remove the pinch clamp. Water will enter into the flask as the air sample cools and decreases in volume.
After the flask has been immersed into the cool water at least 10 minutes, measure the temperature of the cool water bath. It is assumed that the temperature of the cooled air sample in the flask is the same as the cool water outside the flask. Convert the degree Celsius temperature to Kelvin.
Figure 2. Placement of the heated flask into a cool water bath
Keeping the flask inverted (upside down, stopper still pointed downward), raise or lower the flask until the water level inside the flask is equal to the water level in the container (see figure 3). While holding the flask at this level, reattach the pinch clamp to the rubber tubing. Remove the flask from the water, and set it upright on the desk. Remove the clamp and use a graduated cylinder to measure the volume of water that entered the flask as the air sample cooled. Record the volume in ml.
Figure 3. Procedure to equalize the water level inside and outside the flask
Measure the total volume (ml) of the flask by filling it to the top with water, but leave space for the volume occupied by the rubber stopper assembly. Record. Repeat the measurement 3 times!
List the temperatures (K) of the boiling water bath and the total gas volume (ml) obtained at three different runs.
Calculate the average boiling water bath temperature (K) and the average for the total gas volume of the flask.
Repeat the measurement in three different cool water baths, measure the temperature and the volume of water that entered to the flasks.
2.2. Data Analysis
a. Use average temperature for the boiling water baths. Calculate the volume of the air in the flask for each temperature by subtracting the amount of cool water that entered the flask from the average total gas volume.
Volume of cool air = total flask volume - volume of water flask (average)
b. For each sample, calculate the V/T(K) value. These values should be constant, according to the Charles’ law. Any value that is not similar will not be a good data point to use on the graph.
c. Graph the volume-temperature relationship of the gas. (Absolute zero is the theoretical value for the coldest temperature that matter can attain. The value of absolute zero is predicted by extrapolating to the axis where the volume of the gas would decrease to 0 ml.)
d. Boyle’s law
In an experiment, the volume of a specific amount of gas is measured at different volumes, while the temperature is kept constant. The results are summarized in a report sheet table. Determine the P x V product by multiplying the pressure and the volume in each sample. Round the product to give correct number of significant figures.
Mark the vertical axis in equal intervals of mm Hg of pressure. Divide the horizontal axis into equal intervals of ml. For the lowest pressure value, use a pressure that is slightly below the lowest value in the data: the highest value should be just above the highest measured value for pressure. For example, the pressure scale might begin at 600 mm Hg and go up to 1000 mm Hg. The volume values also should begin with a value near the smallest volume obtained. Plot the pressure (mm Hg) of the gas in each reading against the volume (ml). Draw a smooth line through the points obtained from the data. The data points will fall on a slight curve called a hyperbola, not a straight line. Use the graph to discuss the meaning of Boyle’ law.
Pressure / mm Hg / Volume / ml / P x V630 / 32.0
690 / 29.2
726 / 27.8
790 / 25.6
843 / 24.0
914 / 22.2
EXPERIMENT 2
THERMOCHEMISTRY
I. PURPOSE
In this experiment we will investigate the principles of calorimetry and calculate the enthalpy of reaction and neutralization as well as heat capacity.
II. THEORY
Most chemical reactions involve the transfer of energy between the system (the chemical reaction) and the surroundings. One form of energy exchanged is heat, given the symbol q. In some reactions, heat is released when the products are formed (exothermic reaction, q<0) and in other reactions, heat must be added to the system in order for the reaction to proceed (endothermic reaction, q>0). When the reaction is carried out under conditions of constant pressure, the heat exchanged is called the enthalpy of reaction, DHrxn. In SI units, the enthalpy of reaction is expressed in kilojoules per mole of reactant.
The enthalpies of several acid-base reactions and the dissolution of a solid substance will be determined. These enthalpy values will be used to calculate other thermodynamic quantities and the enthalpies for related reactions.
Reactions will be carried out in a calorimeter (a styrofoam cup). A calorie is the amount of energy required to raise the temperature of one gram of water one degree Celsius. In this experiment, you will mix hot and cold water of known temperature and mass. Using the definition of calorie, you will be able to determine the amount of the heat energy that is transferred in bringing the hot and cold water to their final equilibrium temperature, and thereby determine if heat energy is conserved in this process.
Changes in temperature can be measured, from which the energy change (quantity of heat released/absorbed) can be calculated. The enthalpy of the reaction, DHrxn = qsystem, since the experiment is carried out under conditions of constant pressure.
DHrxn = qsystem = -qsurroundings
Where; qsurroundings = Csurr (DTsurr) (msurr).
C = specific heat of the surroundings (J/g.°C)
m = mass of the surroundings (g)
DT = temperature change (°C) = Tfinal - Tinitial
The “surroundings” consist of the calorimeter and its contents. If the calorimeter did not absorb any heat, all of the energy given off by the reaction would raise the temperature of the reaction mixture. However, since some of the heat will be absorbed by the calorimeter, a calorimeter constant, W, must be calculated. Each calorimeter has its own calorimeter constant dependent on its size and composition. This must be experimentally determined.
The reaction will be carried out in water, which has a specific heat capacity of 4.184 J/g.°C. We will assume that the specific heat of the solutions is also 4.184 J/g.°C.
Fig. 1 Experimental Calorimeter
Calculation of DH for the neutralization reactions:
The heat evolved (q) is equal to the sum of the heat absorbed by the reaction mixture and the heat absorbed by the calorimeter. The enthalpy of reaction DH is a molar quantity obtained by dividing the heat evolved, q, by the number of moles of water formed in the neutralization reaction. For an exothermic reaction, DH is negative because heat flows from the system to the surroundings. Calculate DH values for each of the three neutralization reactions to the proper number of significant figures.
Calculation of DH for Related Reactions – Hess’s Law:
Hess’s Law states: The amount of heat generated or absorbed by a chemical reaction is constant regardless of whether the reaction takes place in one or several steps. All chemical reactions which start with the same reactants and end with the same products involve the same net energy change, which is independent of the pathway by which the final state is reached. Enthalpy is a state function.
To illustrate Hess’s Law, the following reactions can be considered:
a) Sodium hydroxide with hydrochloric acid DH1